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Chemistry LibreTexts

15.2: Energy and Chemical Reactions

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  • Learning Objectives

    • Define endothermic and exothermic reactions.
    • Determine whether a reaction is endothermic or exothermic through observations, temperature changes, or an energy diagram.
    • Describe the calculation of heat of reaction using bond energies.

    When physical or chemical changes occur, they are generally accompanied by a transfer of energy. The law of conservation of energy states that in any physical or chemical process, energy is neither created nor destroyed. In other words, the entire energy in the universe is conserved. In order to better understand the energy changes taking place during a reaction, we need to define two parts of the universe, called the system and the surroundings. The system is the specific portion of matter in a given space that is being studied during an experiment or an observation. The surroundings is everything in the universe that is not part of the system. In practical terms for a laboratory chemist, the system is the particular chemicals being reacted, while the surroundings is the immediate vicinity within the room. During most processes, energy is exchanged between the system and the surroundings.

    In the course of an endothermic process, the system gains heat from the surroundings and so the temperature of the surroundings decreases. A chemical reaction or physical change is exothermic if heat is released by the system into the surroundings. Because the surroundings is gaining heat from the system, the temperature of the surroundings increases (Figure \(\PageIndex{1}\)). Exothermic and endothermic reactions can be thought of as having energy as either a product of the reaction or a reactant. Exothermic reactions give off energy, so energy is a product. Endothermic reactions require energy, so energy is a reactant.

    Phase changes, discussed in the previous section 7.3, are also classified in a similar way. The change from gas to liquid (condensation) and liquid to solid (freezing) are exothermic. The change from solid to liquid (melting), and liquid to gas (evaporation and boiling) are endothermic. The exothermic processes release heat to the surroundings while the endothermic processes absorb heat from the surroundings.

    Endothermic reaction: surroundings get cooler and delta H is greater than 0, Exothermic reaction: surroundings get warmer and delta H is less than 0
    Figure \(\PageIndex{1}\). (A) Endothermic reaction. (B) Exothermic reaction.

    When methane gas is combusted, heat is released, making the reaction exothermic. Specifically, the combustion of \(1 \: \text{mol}\) of methane releases 890.4 kilojoules of heat energy. This information can be shown as part of the balanced equation.

    \[\ce{CH_4} \left( g \right) + 2 \ce{O_2} \left( g \right) \rightarrow \ce{CO_2} \left( g \right) + 2 \ce{H_2O} \left( l \right) + 890.4 \: \text{kJ}\]

    The process in the above thermochemical equation can be shown visually in the figure below.

    Figure \(\PageIndex{2}\) (A) As reactants are converted to products in an exothermic reaction, enthalpy is released into the surroundings. The enthalpy change of the reaction is negative. (B) As reactants are converted to products in an endothermic reaction, enthalpy is absorbed from the surroundings. The enthalpy change of the reaction is positive. (CC BY-NC; CK-12)

    In the combustion of methane example, the enthalpy change is negative because heat is being released by the system. Therefore, the overall enthalpy of the system decreases. The heat of reaction is the enthalpy change for a chemical reaction. In the case above, the heat of reaction is \(-890.4 \: \text{kJ}\). The thermochemical reaction can also be written in this way:

    \[\ce{CH_4} \left( g \right) + 2 \ce{O_2} \left( g \right) \rightarrow \ce{CO_2} \left( g \right) + 2 \ce{H_2O} \left( l \right) \: \: \: \: \: \Delta H = -890.4 \: \text{kJ}\]

    Endothermic reactions absorb energy from the surroundings as the reaction occurs. When \(1 \: \text{mol}\) of calcium carbonate decomposes into \(1 \: \text{mol}\) of calcium oxide and \(1 \: \text{mol}\) of carbon dioxide, \(177.8 \: \text{kJ}\) of heat is absorbed. The process is shown visually in the figure above (B). The thermochemical reaction is shown below.

    \[\ce{CaCO_3} \left( s \right) + 177.8 \: \text{kJ} \rightarrow \ce{CaO} \left( s \right) + \ce{CO_2} \left( g \right)\]

    Because the heat is absorbed by the system, the \(177.8 \: \text{kJ}\) is written as a reactant. The heat of reaction is positive for an endothermic reaction.

    Example \(\PageIndex{1}\)

    Label each of the following processes as endothermic or exothermic.

    1. water boiling
    2. gasoline burning
    3. ice forming on a pond


    1. endothermic - you must put a pan of water on the stove and give it heat in order to get water to boil. Because you are adding heat/energy, the reaction is endothermic.
    2. exothermic - when you burn something, it feels hot to you because it is giving off heat into the surroundings.
    3. exothermic - think of ice forming in your freezer instead. You put water into the freezer, which takes heat out of the water, to get it to freeze. Because heat is being pulled out of the water, it is exothermic. Heat is leaving.

    Exercise \(\PageIndex{1}\)

    Label each of the following processes as endothermic or exothermic.

    1. water vapor condensing
    2. gold melting
    Answer (a)
    Answer (b)

    Example \(\PageIndex{2}\)

    Is each chemical reaction exothermic or endothermic?

    1. 2H2(g) + O2(g) → 2H2O(ℓ) + 135 kcal
    2. N2(g) + O2(g) + 45 kcal → 2NO(g)


    1. Because energy is a product, energy is given off by the reaction. Therefore, this reaction is exothermic.
    2. Because energy is a reactant, energy is absorbed by the reaction. Therefore, this reaction is endothermic.

    Exercise \(\PageIndex{2}\)

    Is each chemical reaction exothermic or endothermic?

    1. H2(g) + F2(g) → 2HF (g) + 130 kcal
    2. 2C(s) + H2(g) + 5.3 kcal → C2H2(g)

    a. The energy (130 kcal) is produced, hence the reaction is exothermic

    b. The energy (5.3 kcal) is supplied or absorbed to react, hence, the reaction is endothermic

    Bond Energy

    Atoms bond together to form compounds because in doing so they attain lower energies than they possess as individual atoms. A quantity of energy, equal to the difference between the energies of the bonded atoms and the energies of the separated atoms, is released, usually as heat. That is, the bonded atoms have a lower energy than the individual atoms do. When atoms combine to make a compound, energy is always given off, and the compound has a lower overall energy.

    When a chemical reaction occurs, molecular bonds are broken and other bonds are formed to make different molecules. For example, the bonds of two water molecules are broken to form hydrogen and oxygen.

    \[ 2H_2O \rightarrow 2H_2 + O_2\]

    Energy is always required to break a bond, which is known as bond energy. While the concept may seem simple, bond energy serves a very important purpose in describing the structure and characteristics of a molecule.

    Energy is always required to break a bond. Energy is released when a bond is made.

    Table \(\PageIndex{1}\) Approximate Bond Energies
    Bond Bond Energy (kJ/mol)
    C–H 413
    C–O 358
    C=O* 745
    C–N 305
    C–C 347
    C=C 614
    C≡C 839
    N–H 391
    O—H 467
    H–H 432
    H—Cl 427
    H—I 295
    *C = O(CO2) = 799

    When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to make products. The new arrangement of bonds does not have the same total energy as the bonds in the reactants. Therefore, when chemical reactions occur, there will always be an accompanying energy change. The enthalpy change, for a given reaction can be calculated using the bond energy values from Table \(\PageIndex{1}\).

    In this process, one adds energy to the reaction to break bonds, and extracts energy for the bonds that are formed.

    \[\text{enthalpy change} = \sum (\text{bonds broken}) - \sum (\text{bonds formed})\]

    Example \(\PageIndex{3}\): Generation of Hydrogen Iodide

    What is the enthalpy change for this reaction and is it endothermic or exothermic?

    \[H_2(g)+I_2(g) \rightarrow 2HI(g)\]


    First look at the equation and identify which bonds exist on in the reactants.

    • one H-H bond and
    • one I-I bond

    Now do the same for the products

    • two H-I bonds

    Then identify the bond energies of these bonds from the table above:

    • H-H bonds: 436 kJ/mol
    • I-I bonds: 151 kJ/mol

    The sum of enthalpies on the reaction side is:

    436 kJ/mole + 151 kJ/mole = 587 kJ/mol.

    This is how much energy is needed to break the bonds on the reactant side. Then we look at the bond formation which is on the product side:

    • 2 mol H-I bonds: 297 kJ/mol

    The sum of enthalpies on the product side is:

    2 x 297 kJ/mol= 594 kJ/mol

    This is how much energy is released when the bonds on the product side are formed. The net change of the reaction is therefore

    587 kJ/mol -594 kJ/mol= -7 kJ/mol.

    Since this is negative, the reaction is exothermic.

    Exercise \(\PageIndex{3}\): Decomposition of Water

    Using the bond energies given in the chart above, find the enthalpy change for the thermal decomposition of water:

    \[ 2H_2O (g) \rightarrow 2H_2 + O_2 (g) \]

    Is the reaction written above exothermic or endothermic? Explain.


    The enthalpy change deals with breaking two mole of O-H bonds and the formation of 1 mole of O-O bonds and two moles of H-H bonds (Table \(\PageIndex{1}\)).

    • The sum of the energies required to break the bonds on the reactants side is 4 x 460 kJ/mol = 1840 kJ/mol.
    • The sum of the energies released to form the bonds on the products side is
      • 2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol
      • 1 moles of O=O bond = 1 x 498.7 kJ/mil = 498.7 kJ/mol

    which is an output (released) energy = 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol.

    Total energy difference is 1840 kJ/mol – 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ of heat is needed to be supplied to carry out this reaction.

    This reaction is endothermic since it requires energy in order to create bonds.


    • Chemical processes are labeled as exothermic or endothermic based on whether they give off or absorb energy, respectively.
    • Atoms are held together by a certain amount of energy called bond energy.
    • Energy is released to generate bonds, which is why the enthalpy change for breaking bonds is positive. Energy is required to break bonds. Atoms are much happier when they are "married" and release energy because it is easier and more stable to be in a relationship (e.g., to generate octet electronic configurations). The enthalpy change is negative because the system is releasing energy when forming bond.

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